Chem 110 Lab
Clark College
Observing, Writing, and Predicting Chemical Reactions
Name: ___________________________________
Partner’s Name: ___________________________________
Introduction
As someone taking a chemistry class, you will need to understand chemical formulas, chemical
reactions, and how species interact with one another. All of these things apply to you in your
everyday life – as you use shampoos, soaps, and conditioners, as you eat food and digest them in
your body. As you inhale and exhale – air and contaminants. Understanding the nature of
chemicals requires the ability to write a compound’s correct chemical formula and the ability to
predict some of the simple chemical reactions that a compound may undergo.
LEARNING OBJECTIVES After completing this experiment, you should feel comfortable with:
I.
•
Working with solutions.
•
Proper waste disposal.
•
Making detailed observations of chemical reactions
•
Writing and balancing single and double replacement reactions, including predicting
products.
Writing Formulas for Compounds
When atoms of different elements combine, new substances called compounds are formed. An
equation is a chemist’s method of quantitatively describing how elements combine to form
compounds. Since there are more elements (over 100) than letters of the alphabet (26), many of the
elements are represented by two letter symbols. Because some elements are represented by a single
letter, it is important to write the first letter as a capital letter and the second one as a small letter (as
examples we write Sn, Bi, Cu, rather than SN, BI, CU).
Numbers in chemical formulas have different meanings depending upon their position. Consider,
for example, Ca3(PO4)2. The “3” and “4” which directly follow the atomic symbol as subscripts refer
only to that symbol, indicating 3 atoms of Ca and 4 atoms of O, respectively. The “2” following the
parenthesis multiplies everything inside the parentheses by 2.
Therefore in the formula Ca3(PO4)2, you have the following:
Ca3(PO4)2
3 Ca atoms, 2 P atoms, and 8 O atoms.
Numbers that appear as superscripts to the right of the symbol indicate the charge on the symbol:
for example Ca+2 or PO4-3. Because this compound consists of ions, the subscripts “3” and “2” also
indicate that you have 3 Ca+2 ions and 2 PO4-3 ions. These charges are left out of the final
balanced formula of the neutral compound.
The Periodic Table allows you to determine the valence electron number for all Main Group (Group
A, Columns 1-8) elements. The number of valence (outer shell) electrons is indicated by the Roman
Numeral above each group. All Group IA atoms have 1 valence electron, all Group IIA atoms have
2 valence electrons, etc.
Chemical Reactions
Spring 2010 AEM
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Rules for Writing Formulas for Ionic Compounds
1. Atoms containing less than 4 valence electrons tend to lose all their valence electrons, forming
positively charged ions (Na+, Mg+2,Al+3).
2. Atoms containing more than 4 valence electrons tend to gain electrons until the outer shell
contains an octet of electrons, forming negatively charged ions (N-3, O-2, F-1).
Using these two rules and following the sequence of steps we can write the correct formulas for
ionic compounds:
1. Write the symbols for the atoms in the compound. Positive ions (usually metals) appear first,
followed by the negatively charged ion (nonmetals) or polyatomic ion.
2. Write the charge (obtained by using the Periodic Table and the Octet Rule) for each atom (or
polyatomic ion) above and to the right of the symbol.
3. If charges are different from each other, use the numbers of the charges only, drop the sign,
exchange them and write them as subscripts, omitting subscripts that have a subscript of 1.
4. Check your answer by multiplying charge by subscript for each ion and adding; the total charge
should equal zero.
Example: Write the formula for aluminum oxide
1. Al O
Write the positive ion (atom) first, negative ion (atom) last.
+3
-2
Charge of Al = +3, charge of O = -2.
2. Al O
3. Al2O3
Exchange charges and write them as subscripts.
4. For Al (+3) x (2) = +6; for O (-2) x (3) = -6; the compound is then (+6) +(-6) = 0
I.
Chemical Equations
A. Balancing a Chemical Equation
In balancing a chemical equation you must always observe two basic principles. These are:
1. Atoms are neither created nor destroyed in a chemical reaction.
Example: H2 + Cl2
HCl
Correct formulas and product, but atoms not balanced
2. Subscripts of correctly written formulas for compounds or molecules may not be changed to
balance atoms.
Example: H2 + Cl2
H2Cl2
Atoms are balanced but equation is incorrect since subscripts changed.
Correct:
H2 + Cl2
2 HCl
To balance a chemical equation, balance one atom (or polyatomic ion) at a time, using only
coefficients leaving any that occur as elements until last.
Examples:
Al + CuSO4
* 2 Al + 3 CuSO4
Chemical Reactions
Al2(SO4)3 + Cu
Al2(SO4)3 + 3 Cu
Spring 2010 AEM
all species balanced
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MgI2 + H3PO4
Mg3(PO4)2 + HI
* 3 MgI2 + 2 H3PO4
Mg3(PO4)2 + 6 HI
all balanced
Practice: Determine the coefficients needed to balance each equation.
a. _____H3PO4 + _____Ca(OH)2
_____HOH
+
_____Ca3(PO4)2
b. _____Mg
+ _____N2
_____Mg3N2
c. _____KOH
+ _____H2SO4
_____K2SO4 +
_____HOH
d. _____Sn
+ _____CuF
_____SnF2
_____Cu
+
B. Writing (and Balancing) Chemical Equations
Before you can balance a chemical equation, you must first write the formulas for the reactants and
the products. Use the steps below to obtain a correct chemical equation.
1. On the left side of the arrow write the correct formulas for the reactants.
2. Use the nature of each reactant (element, compound) to determine the type of reaction
(decomposition, combination, single replacement, double replacement) that will occur, and
write the correct formula for the products. **Note that H2, N2, O2, F2 Cl2, Br2, and I2 occur as
diatomic molecules in nature.
3. After correct formulas for reactants and products are written using subscripts, balance the
equations with the previously described procedure. Use only coefficients so that you have
equal numbers of each type of atom (or polyatomic ion) in the reactant(s) and in the product(s).
II. Types of Chemical Reactions
The two important types of chemical reactions are:
A. Single Replacement (SR)
B. Double Replacement (DR)
These are described in the following sections.
A. Single Replacement:
One metal replaces either hydrogen or another metal from a compound in solution. Single
replacement (SR) reactions are always redox reactions.
Recognition:
Anytime a free element and a compound are reactants, look for a single replacement reaction,
in which the free element replaces an element from the compound, yielding a different free
element and a different compound. The charge on a free element is zero.
Chemical Reactions
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Examples:
Metal replaces metal ion
Mg(s)
+ CuSO4(aq)
MgSO4(aq)
+
Cu(s)
Metal replaces hydrogen ion from acid
6K(s)
+
2H3PO4(aq)
2K3PO4(aq)
+
H2(g)
Instructor Demonstrations:
The instructor will perform the first reaction for you while you all observe. Reactions 1, 2, and 3
will be started while you watch, and then you will have an opportunity to observe the results
individually during the lab period. Describe the appearance of each substance before the
reaction. Allow the reaction to proceed for at least 10 minutes, and make final observations.
before you dispose of the mixture in the appropriate container. Write balanced equations for each
reaction that occurred. In the cases in which there is no reaction, write “NR.”
1. Sodium metal + water
observations before:
observations during & after reaction:
balanced equation:
→
_____Na (s) + _____H2O (l)
2. Copper metal + water (note, if the copper (II) ion is formed in solution the solution will be
blue). If no reaction is observed, for the balanced equation write NR
observations before:
observations during & after reaction:
balanced equation: _____Cu (s) + _____H2O (l)
→
3. Copper solid + aqueous silver nitrate (note: if the copper (II) ion is formed in solution, the
solution will turn blue, if no reaction is observed, in the balanced equation, put NR)
observations before:
observations during & after reaction:
balanced equation:
Chemical Reactions
_____Cu (s) + _____AgNO3 (aq)
Spring 2010 AEM
→
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B. Double Replacement:
Positive and negative ions exchange partners. Double Replacement (DR) reactions are NOT
redox reactions. Ions keep the same charge from reactants to products.
Recognition:
Two compounds must be present, and the exchange in ions must produce of one of these:
1. a liquid (usually water).
2. a gas (ex: H2S or CO2)
3. a precipitate – or solid (water insoluble solid). Use the solubility rules.
Examples:
1. Formation of a non-ionizable substance (water)
2 NH4OH(aq) +
H2SO4(aq)
(NH4)2SO4(aq) + 2 HOH(l)
Na2S(aq)
H2S(g)
2. Formation of a gas
a.
H2SO4(aq)
+
+
Na2SO4(aq)
Formation of a precipitate (refer to solubility rules listed below)
a.
AgNO3(aq)
b.
c.
KCl(aq)
+
HBr(aq)
AgBr (s)
+
HNO3(aq)
3 Na2CO3(aq) +
2 AlCl3(aq)
6NaCl(aq)
+
Al2(CO3)3 (s)
+
NaNO3(aq)
no reaction, all products remain aqueous
Solubility Rules
Compounds with…
alkali metals (Li+, Na+, etc.) and
ammonia (NH4+)
Solubility
Exceptions
soluble (aq)
None
nitrates (NO3-)
soluble (aq)
acetates (C2H3O2-)
soluble (aq)
chlorates (ClO3-)
chlorides, bromides, iodides
(Cl-, Br-, I-)
soluble (aq)
sulfates, SO42-
soluble (aq)
hydroxides, OH-
insoluble (s)
carbonates, CO32-
insoluble (s)
phosphates, PO43-
insoluble (s)
Chemical Reactions
soluble (aq)
Spring 2010 AEM
None
None
None
halides of Ag+, Pb2+, Hg22+
are insoluble (s)
sulfates of Ba2+, Sr2+, Pb2+, Hg22+
are insoluble (s)
hydroxides of alkali metals,
NH4+,Ba2+, Ca2+, Sr2+
are soluble (aq)
carbonates of alkali metals and NH4+
are soluble( aq)
phosphates of alkali metals and NH4+
are soluble (aq)
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Experimental:
You have several aqueous solutions and will perform reactions between each possible pair of
solutions. From the reagents bench, obtain a plastic sheet marked like the grid. In each square of
the grid, place two-three drops of each appropriate solutions in their grid box. *** (note if you are
pregnant or thinking of becoming pregnant, please wear gloves for this experiment). Make
your observations of what the solutions look like before mixing them.
In the other grid boxes you will mix together the two solutions as indicated. For example, the
compound on the left of the row tells you what solution to add in each box ACROSS the row
and the compound on the top of the columns tells you what to add DOWN a column. This
means that when mixing, each grid box will contain 2 chemicals that are mixed. (use 2 drops of
each for mixing). Make your observations. After you perform your solution binary mixing you
should wipe off the solution residue with a paper towel. Place the paper towel in the trash, and
then you can wash off your plastic reaction grid with soap and water in the sink.
Each square will contain two solutions, and you should take care to place the second drop on
top of the first drop without contaminating the droppers. In the data table, describe your
observations. Look for formation of a solid precipitate (sometimes this looks like cloudiness, or
the presence of a powdery substance suspended in the water). If no reaction occurs write NR.
Results:
For each pure solution, record your initial observations, for each of your binary mixture of
solutions, record your observations. Be sure to take care with your observations and be as
descriptive as possible!
Reflection assignment:
For each reaction that produces a solid precipitate, write the balanced equation, including phase
labels.
Chemical Reactions
Spring 2010 AEM
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